Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

Chemistry > EL 8: Spectacular Metals > Flashcards

EL 8: Spectacular Metals Flashcards

You may prefer our related Brainscape-certified flashcards:
Periodic Table flashcards Chemistry 101 flashcards AP® Chemistry flashcards Organic Chemistry 101 flashcards
1
Q

What is the periodic table arranged into?

A

Periods (rows) and groups (columns)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

What do all the elements in the same period have?

A

The same number of electron shells

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

What do all the elements within the same group have?

A

The same number of electrons in their outer shells (The group number tells you the number of electrons on the outermost electron shell)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

What typically happens to the properties of elements as you go down the group?

A

The relativity changes, it usually increases or decreases down the group

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

What do periods 2 and 3 show similar trends in?

A

Melting point

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

What does melting point have to do with bonding and structure?

A

In order for a substance to melt, it has to overcome the forces between its atoms or molecules so these trends are linked to changes in structure and bond length

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

GRAPH

A

For the group 2 metals (Li and Be) and the group 3 metals (Na, Mg and Al), the melting points increase across the period because the metal-metal bonds get stronger.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

Why do the metal bonds get stronger?

A

The bonds get stronger because the metal ions have an increasing number of delocalised electrons and a decreasing ionic radius (the ionic radius decreases because the electrons are pulled in closer as the number of protons in the nucleus increases).

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

What does stronger metal bonding lead to?

A

This leads to a higher charge density, which attracts the ions together more strongly

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

Why is there a peak in the graph at carbon and silicone?

A

Carbon and silicone are giant covalent substances, they have strong covalent bonds linking all their atoms together, therefore a lot of energy is required to break these bonds (carbon and silicon have the highest melting points in their periods)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

Why do the group 2 elements (N, O and F) and the group 3 elements (P, S and Cl) have low melting points?

A

All of these elements are simple molecular covalent substances, their melting points depend upon the strength of the intermolecular forces between individual molecules, because they have low melting points this means that they have weak intermolecular forces and are therefore easily broken (low melting points)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Explain why there is a slight increase in melting point of the substance Sulphur between the elements phosphorus and chlorine

A

Sulphur is the biggest molecule is period 3, therefore it has more atoms (More atoms mean stronger intermolecular forces), so its got a higher melting point than phosphorus and chlorine.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

Explain why the atoms Neon (Ne) and Argon (Ar) have the lowest melting points in the graph

A

Neon and argon are noble gases and have the lowest melting point because they exist as individual atoms (monatomic) which results in very weak intermolecular forces and therefore low melting points.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

What is the term given to when electrons have been removed from an atom or molecule?

A

Ionised

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

what is the energy you need to remove the first electron to ionise it?

A

First ionisation enthalpy/energy

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

Give an example equation

A

O(g) ——> O+(g) + e-

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

What is the relationship between how low the ionisation energy is and the ability to easily remove an atom?

A

The lower the ionisation enthalpy, the easier it is to remove an out electron and form an ion

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

What are the 3 main things that affect first ionisation energy?

A

Atomic radius
Nuclear Charge
Electron shielding

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

How does Atomic radius affect First ionisation energy?

A

The further the outer shell electrons are from the positive nucleus, the less they’ll be attracted towards the nucleus, so the first ionisation energy will be lower

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

How does nuclear charge affect First ionisation energy?

A

This is the positive charge on the nucleus, caused by the presence of protons, the more protons there are in the nucleus, the more it’ll attract the outer electrons - it’ll be harder to remove the outer electrons, so the ionisation enthalpy will be higher

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

How does Electron shielding affect First ionisation energy?

A

The inner electron shells shield the outer shell electrons from the attractive force of the nucleus, because more inner shells mean greater electronic shielding, the ionisation enthalpy will be lower.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

What groups does the S-block contain?

A

Groups 1 and 2

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

What happens to first ionisation energy/enthalpy as you go down a group?

A

First ionisation energy decreases as you go down the group, the outer electrons are in shells further from the nucleus so they’re attracted to the nucleus less. The amount of shielding increase as you go down the group because there are more filled inner shells, this means less nuclear attraction for the outer shell electrons so they’re easier to remove.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

What happens to first ionisation energy/enthalpy as you go across a period?

A

First ionisation energy increases as you go across a period. This is because the number of protons is increasing so the outer electrons are more strongly attracted to the nucleus. Also the outer shell electrons are all roughly the same energy level so the level of electron shielding is the same. Therefore it is more difficult to remove an outer electron and the first ionisation enthalpy is higher.

25
Q

Why do S-block metals have low ionisation enthalpies?

A

The metals in the S-block have relatively low nuclear charges, so they have low first ionisation energies compared to other elements in the same period - this means they lose their outer electrons easily as there is less attraction between the nucleus and the outer electrons (makes the s-block metals reactive).

26
Q

Give a use of Beryllium oxide, BeO

A

Nose-cones of rockets because it has a high melting point

27
Q

Give a use of magnesium hydroxide, Mg(OH)2

A

Some indigestion tablets

28
Q

Give a use of Calcium Carbonate, CaCO3

A

Agriculture to neutralise acid soils because its insoluble and reacts with acid

29
Q

Give a use of Calcium sulphate hydrated, CaSO4

A

Plaster to keep broken limbs in place because its insoluble

30
Q

Give a use of strontium carbonate, SrCO3

A

Fireworks because a red colour is given off when heated

31
Q

Give a use of Barium Sulphate, BaSO4

A

Barium meal to allow x-raying of digestive tract because its opaque to x-rays and is insoluble

32
Q

Give a use of Radium chloride RaCl2

A

Treat certain cancers because its radioactive

33
Q

What is the trend with group 1 and 2 as you go down the group?

A

Elements become more metallic, they more readily form cations in ionic compounds, for this reason the most reactive metals in groups 1 and 2 are found at the bottom of the group

34
Q

What is the trend with group 1 and 2 as you go across the period?

A

Elements become less metallic as you go across the period from left to right, for this reason group 1 metals are more reactive than group 2 metals in the same period

35
Q

What happens to atomic radius as you go down a group?

A

Increases, this is because the greater the atomic number the more electrons there are and these go into shells increasingly further from the nucleus.

36
Q

What happens to atomic radius as you go across a period?

A

Decreases, this is because atoms of group 2 are smaller than the equivalent of the group 1 atom, the extra proton exerts a greater attraction to the electrons = smaller radii

37
Q

Reactivity of group 2 compounds with oxygen

A

Reactivity increases as you go down the group, due to the ease of cation formation and react with increasing vigour down the group. Atoms are larger as you down the group (increasing atomic radius) and there is more electron shielding so its easier to remove the outer electron.

38
Q

What do group 2 elements react with oxygen to produce?

A

A metal oxide

39
Q

What happens when group 2 elements react?

A

They form positive ions with a charge of +2, this is because group 2 atoms contain 2 electrons in their outer shell

40
Q

What happens when group 2 elements react with water?

A

They react to produce hydrogen gas and a metal hydroxide. Reactivity increases as you go down the group, due to the ease of cation formation and react with increasing vigour down the group. Atoms are larger as you down the group (increasing atomic radius) and there is more electron shielding so its easier to remove the outer electron. (They get increasingly reactive down the group because the outermost electrons are further from the nucleus and so more easily lost.

41
Q

What is the general formula for Group 2 carbonates?

A

MCO3

42
Q

Define “thermal decomposition”

A

The breaking down of substances using heat

43
Q

What two products do group 2 carbonates decompose to form?

A

An oxide and carbon dioxide

44
Q

As you go down the group of group 2 carbonates, do they become more or less difficult to thermally decompose?

A

More difficult

45
Q

Why do group 2 carbonates become more difficult to decompose further down the group?

A

They are more thermally stable and therefore take more heat to break down

46
Q

Define “charge density”

A

The charge of an ion relative to its volume

47
Q

What is the relation of an ion’s charge density to its size?

A

The smaller an ion, the higher its charge density

48
Q

How do group 2 metal cations affect the carbonate anions?

A

They draw electrons on the carbonate anion towards itself, polarising the negative charge cloud around the anion, distorting it and making it less stable.

49
Q

Why are group 2 carbonates more thermally stable the further down the group?

A
  • The further down the group, the larger the cations
  • Large cations cause less distortion than small cations as they have a lower charge density
  • Therefore the further down the group, the more stable the carbonate anions
50
Q

What is the general formula for Group 2 oxides?

A

MO

51
Q

What is the general formula for Group 2 hydroxides?

A

M(OH)2

52
Q

What do the oxides of Group 2 metals react with water to form? Give the formula.

A

Metal hydroxides

MO + H2O ====> M+ + OH-

53
Q

What makes the solutions formed from the reactions of Group 2 metal oxides with water strongly alkaline?

A

Because hydroxide ions dissolve in the water

54
Q

How does magnesium oxide react with water?

A

Slowly, forming a not very soluble hydroxide

55
Q

What is the relation between going down the Group 2 group and the strength of the alkaline solutions formed by their oxides and why?

A

The further down the group, the ore strongly alkaline solutions formed, because the solubility of the hydroxides increase.

56
Q

Why do group 2 metal oxides and hydroxides neutralise dilute acids?

A

They are bases

57
Q

What do group 2 metal oxides and hydroxides form on reacting with dilute acids? Give the formulas for a reaction with hydrochloric acid

A

Solutions of the corresponding salts
MO(s) + 2HCl (aq) ====> MCl2 (aq) + H2O (l)
M(OH)2(s) + 2HCl (aq) ====> MCl2 (aq) + 2H2O (l)

58
Q

What is the neutralising effect of the reaction of group 2 oxides/hydroxides with acid used for?

A

By farmers for putting lime (calcium hydroxide) on their fields to neutralise soil acidity.

59
Q

How does solubility of Group 2 metal carbonates change down the group?

A

It decreases

Chemistry (11 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions